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| Chem 110 |
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Fall 2003 |
| Lecture Notes::Lec 6_8 September |
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| PREVIOUS |
Even though energy can neither be created nor destroyed, the form of the energy can change.
For our purposes we need to consider three forms of energy:
Another term we must be aware of is Heat. Heat is a general catch-all for energy in transit. Notice that heat always moves spontaneously from hotter to cooler objects.
Recall the temperature is a measure of the average KE of the particles in a system. It can thus serve as an indicator of how readily the energy in a system can be transferred. However, temperature alone does not tell us how much energy is available in a system. We need to know not only the temperature, but also the heat capacity of the system to know how much energy it holds.
Demonstration: Boil water using a gas burner. Note that the energy (heat) released by the burner flame results from the breaking and making bonds between atoms. The energy differences in these bonds is released as KE (the molecules fly off at higher speeds) and radiation. This KE is then transferred to the screen and beaker etc.
The First Law tells us about the energy gained or lost during a process, but not whether that process will occur. So how do we predict which processes are likely to happen?
First we need to define the term spontaneous. In chemistry spontaneous means that a process will occur without outside influence in the direction written. In terms of equilibria we can say that a spontaneous reaction is one where the products are favored over the reactants (there is at least a slight excess).
{Note that saying a reaction is spontaneous says nothing about how fast the reaction or process occurs, just that it will occur! In other words the kinetics is not related to thermodynamics of a reaction. Thermodynamic functions are intrinsic to the overall reaction, and can't be changed without changing the reactants and products (under the same conditions) - they are pathway independent. Kinetics on the other hand can change drastically with path (such as with catalysts).}
So what determines whether a process is spontaneous? Two components contribute under common laboratory and biological conditions:
More generally we can say that the entropy of the Universe increases in a spontaneous process. This is enshrined as the second law of thermodynamics:
The Second Law of Thermodynamics: The Entropy to the Universe increases in any spontaneous process.
Note that while the First Law says the Energy of the universe is constant, the entropy is not. Entropy has sometimes been referred to as "times arrow" because it is what gives the Universe a direction - we can only go forward in time.
So what is entropy? Entropy is often described as a measure of disorder (randomness).
Examples: Opening a small vail of gas in a room allows the gas to disperse throughout the room. The gas has now become more disordered - it is much harder to localize the positions of the gas particles in the larger space. The gas in the large space has more entropy. It is much easier to describe the positions of the ions in a small crystal than when they are dissolved into a large volume of water. The ions dissolved in water have more entropy.
Aside: What about processes which lead to more ordered systems? For example babies eat some pretty disordered looking stuff (pureed food), but grow into highly organized beings. Is the Second Law being violated? No. Although the material incorporated into the baby is more ordered than its food, most of the food was transformed into carbon dioxide and water which dispersed over the world's atmosphere - it has become MUCH more disorganized. As a result the Universe has increased in entropy. Similarly, though even less obviously, plants organize small molecules at the expense, ultimately, of the Sun's dispersal over the Universe.
Another way of describing entropy is to describe the number of arrangements available to a system. (Example on p795 of text.)
If the entropy increase of the Universe is the ultimate driving force for reactions, what's going on? There must be a relationship between DH in the system and the entropy of the Universe. Apparently the transfer of heat affects entropy.
Let's look at how this might occur. When heat is transferred to the surroundings the result is an increase in the temperature, T, of the surroundings. As a result the KE and thus motion of the particles is increased, which in turn means their distribution is harder to describe - it has become more random. Entropy in the surroundings has increased! So heat and entropy are indeed connected when we look at the entire system.
We must consider the system and its surroundings to determine the overall change in entropy and thus the spontaneity of a given process. In most cases we need only analyze the local surroundings to understand what is happening, though ultimately it is the entire Universe which determines spontaneity.
If we look at an exothermic reaction as an example, the reaction mixture loses bond energy as enthalpy, while the surroundings gain heat which increases its entropy. The entropy increase from the enthalpy of the reaction results in a net increase in the entropy of the Universe and thus favors the reaction. Because of this enthalpic entropy increase most exothermic reactions are favored, however in some reactions the entropy of the reaction decreases sufficiently to make them non-spontaneous.
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© R A Paselk
Last modified 9 September 2003
